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CHEMISTRY
CHAPTER
Table Of Contents
6
Matter and Change
Chapter 6: The Periodic Table and
Periodic Law
Section 6.1
Development of the Modern
Periodic Table
Section 6.2
Classification of the Elements
Section 6.3
Periodic Trends
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SECTION
6.1
Development of the Modern Periodic Table
• Trace the development of the periodic table.
• Identify key features of the periodic table.
atomic number: the number of protons in an atom
The periodic table evolved over time as
scientists discovered more useful ways to
compare and organize the elements.
SECTION
6.1
Exit
Development of the Modern Periodic Table
periodic law
group
period
representative elements
transition elements
metal
alkali metals
alkaline earth metals
transition metal
inner transition metal
lanthanide series
actinide series
nonmetals
halogen
noble gas
metalloid
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SECTION
6.1
Development of the Modern Periodic Table
SECTION
6.1
Development of the Modern Periodic Table
Development of the Periodic Table (cont.)
Development of the Periodic Table (cont.)
• In the 1700s, Lavoisier compiled a list of all
the known elements of the time.
• Newlands noticed when the elements
were arranged by increasing atomic
mass, their properties repeated every
eighth element.
• The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
• John Newlands proposed an arrangement
where elements were ordered by increasing
atomic mass.
SECTION
6.1
Development of the Modern Periodic Table
SECTION
6.1
Development of the Modern Periodic Table
Development of the Periodic Table (cont.)
The Modern Periodic Table
• Meyer and Mendeleev both demonstrated
a connection between atomic mass and
elemental properties.
• Moseley rearranged the table by increasing
atomic number, and resulted in a clear
periodic pattern.
• The modern periodic table contains boxes
which contain the element's name, symbol,
atomic number, and atomic mass.
• Periodic repetition of chemical and physical
properties of the elements when they are
arranged by increasing atomic number is
called periodic law.
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SECTION
6.1
Development of the Modern Periodic Table
SECTION
6.1
Development of the Modern Periodic Table
The Modern Periodic Table (cont.)
The Modern Periodic Table (cont.)
• Columns of elements are called groups.
• Elements are classified as metals, non-metals,
and metalloids.
• Metals are elements that are generally shiny
when smooth and clean, solid at room
temperature, and good conductors of heat and
electricity.
• Rows of elements are called periods.
• Elements in groups 1,2, and 13-18 possess a
wide variety of chemical and physical
properties and are called the representative
elements.
• Elements in groups 3-12 are known as the
transition metals.
SECTION
6.1
Development of the Modern Periodic Table
• Alkali metals are all the elements in group 1
except hydrogen, and are very reactive.
• Alkaline earth metals are in group 2, and are
also highly reactive.
SECTION
6.1
Development of the Modern Periodic Table
The Modern Periodic Table (cont.)
The Modern Periodic Table (cont.)
• The transition elements are divided into
transition metals and inner transition
metals.
• Non-metals are elements that are
generally gases or brittle, dull-looking
solids, and poor conductors of heat and
electricity.
• The two sets of inner transition metals are
called the lanthanide series and actinide
series and are located at the bottom of the
periodic table.
• Group 17 is composed of highly reactive
elements called halogens.
• Group 18 gases are extremely unreactive and
commonly called noble gases.
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SECTION
6.1
Development of the Modern Periodic Table
SECTION
6.2
Classification of the Elements
The Modern Periodic Table (cont.)
• Metalloids have physical and chemical
properties of both metals and non-metals,
such as silicon and germanium.
• Explain why elements in
the same group have
similar properties.
• Identify the four blocks
of the periodic table
based on their electron
configuration.
valence electron:
electron in an atom's
outermost orbitals;
determines the chemical
properties of an atom
Elements are organized into different
blocks in the periodic table according
to their electron configurations.
SECTION
6.2
Classification of the Elements
SECTION
6.2
Classification of the Elements
Organizing the Elements by Electron
Configuration
Organizing the Elements by Electron
Configuration (cont.)
• Recall electrons in the highest principal energy
level are called valence electrons.
• All group 1 elements have one valence
electron.
• The energy level of an element’s valence
electrons indicates the period on the
periodic table in which it is found.
• The number of valence electrons for
elements in groups 13-18 is ten less than
their group number.
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SECTION
6.2
Classification of the Elements
SECTION
6.2
Classification of the Elements
The s-, p-, d-, and f-Block Elements
The s-, p-, d-, and f-Block Elements (cont.)
• The shape of the periodic table becomes clear
if it is divided into blocks representing the
atom’s energy sublevel being filled with
valence electrons.
• s-block elements consist of group 1 and 2,
and the element helium.
• Group 1 elements have a partially filled s
orbital with one electron.
• Group 2 elements have a completely filled s
orbital with two electrons.
SECTION
6.2
Classification of the Elements
SECTION
6.2
Classification of the Elements
The s-, p-, d-, and f-Block Elements (cont.)
The s-, p-, d-, and f-Block Elements (cont.)
• After the s-orbital is filled, valence
electrons occupy the p-orbital.
• Groups 13-18 contain elements with
completely or partially filled p orbitals.
• The d-block contains the transition metals
and is the largest block.
• There are exceptions, but d-block elements
usually have filled outermost s orbital, and
filled or partially filled d orbital.
• The five d orbitals can hold 10 electrons, so
the d-block spans ten groups on the periodic
table.
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SECTION
6.2
Classification of the Elements
SECTION
6.3
Periodic Trends
The s-, p-, d-, and f-Block Elements (cont.)
• The f-block contains the inner transition
metals.
• f-block elements have filled or partially filled
outermost s orbitals and filled or partially filled
4f and 5f orbitals.
• Compare period and
group trends of several
properties.
• Relate period and group
trends in atomic radii to
electron configuration.
ion
ionization energy
octet rule
electronegativity
Trends among elements in the periodic
table include their size and their ability to
lose or attract electrons
• The 7 f orbitals hold 14 electrons, and the
inner transition metals span 14 groups.
SECTION
6.3
Periodic Trends
principal energy level:
the major energy level of
an atom
SECTION
6.3
Periodic Trends
Atomic Radius
Atomic Radius (cont.)
• Atomic size is a periodic trend influenced
by electron configuration.
• For elements that occur as molecules, the
atomic radius is half the distance between
nuclei of identical atoms.
• For metals, atomic radius is half the distance
between adjacent nuclei in a crystal of the
element.
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SECTION
Periodic Trends
6.3
SECTION
6.3
Periodic Trends
Atomic Radius (cont.)
Atomic Radius (cont.)
• There is a general decrease in atomic
radius from left to right, caused by
increasing positive charge in the nucleus.
• Atomic radius generally increases as you
move down a group.
• The outermost orbital size increases down a
group, making the atom larger.
• Valence electrons are not shielded from the
increasing nuclear charge because no
additional electrons come between the
nucleus and the valence electrons.
SECTION
Periodic Trends
6.3
SECTION
6.3
Periodic Trends
Ionic Radius
Ionic Radius (cont.)
• An ion is an atom or bonded group of atoms
with a positive or negative charge.
• When atoms lose electrons and form positively
charged ions, they always become smaller for
two reasons:
• When atoms gain electrons, they can
become larger, because the addition of an
electron increases electrostatic repulsion.
1. The loss of a valence electron can leave an empty
outer orbital resulting in a small radius.
2. Electrostatic repulsion decreases allowing the
electrons to be pulled closer to the radius.
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SECTION
6.3
Periodic Trends
SECTION
6.3
Periodic Trends
Ionic Radius (cont.)
Ionic Radius (cont.)
• The ionic radii of positive ions generally
decrease from left to right.
• The ionic radii of negative ions generally
decrease from left to right, beginning with group
15 or 16.
• Both positive and negative ions increase in
size moving down a group.
SECTION
6.3
Periodic Trends
SECTION
6.3
Periodic Trends
Ionization Energy
Ionization Energy (cont.)
• Ionization energy is defined as the energy
required to remove an electron from a
gaseous atom.
• Removing the second electron requires more
energy, and is called the second ionization
energy.
• The energy required to remove the first
electron is called the first ionization energy.
• Each successive ionization requires more
energy, but it is not a steady increase.
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SECTION
6.3
Periodic Trends
Ionization Energy (cont.)
SECTION
6.3
Periodic Trends
Ionization Energy (cont.)
• The ionization at which the large increase
in energy occurs is related to the number of
valence electrons.
• First ionization energy increases from left to
right across a period.
• First ionization energy decreases down a
group because atomic size increases and
less energy is required to remove an electron
farther from the nucleus.
SECTION
6.3
Periodic Trends
SECTION
6.3
Periodic Trends
Ionization Energy (cont.)
Ionization Energy (cont.)
• The octet rule states that atoms tend to
gain, lose or share electrons in order to
acquire a full set of eight valence electrons.
• The electronegativity of an element indicates
its relative ability to attract electrons in a
chemical bond.
• The octet rule is useful for predicting what
types of ions an element is likely to form.
• Electronegativity decreases down a group
and increases left to right across a period.
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SECTION
6.3
Periodic Trends
Ionization Energy (cont.)
10